The materials to prepare students for practical lessons of inorganic chemistry

LESSON 19.

THEME. Phosphorus and its compounds. p-elements of the VA group. Sub-group of Arsenic

 

PHOSPHORUS

Phosphorus is a chemical element with symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus as a mineral is almost always present in its maximally oxidised state, as inorganic phosphate rocks. Elemental phosphorus exists in two major forms—white phosphorus and red phosphorus—but due to its high reactivity, phosphorus is never found as a free element on Earth.

The first form of elemental phosphorus to be produced (white phosphorus, in 1669) emits a faint glow upon exposure to oxygen – hence its name given from Greek mythology, Φωσφόρος meaning “light-bearer” (Latin Lucifer), referring to the “Morning Star”, the planet Venus. The term “phosphorescence”, meaning glow after illumination, originally derives from this property of phosphorus, although this word has since been used for a different physical process that produces a glow. The glow of phosphorus itself originates from oxidation of the white (but not red) phosphorus— a process now termed chemiluminescence.

The vast majority of phosphorus compounds are consumed as fertilisers. Other applications include the role of organophosphorus compounds in detergents, pesticides and nerve agents, and matches.

Phosphorus is essential for life. As phosphate, it is a component of DNA, RNA, ATP, and also the phospholipids that form all cell membranes. Demonstrating the link between phosphorus and life, elemental phosphorus was historically first isolated from human urine, and bone ash was an important early phosphate source. Phosphate minerals are fossils. Low phosphate levels are an important limit to growth in some aquatic systems. The chief commercial use of phosphorus compounds for production of fertilisers is due to the need to replace the phosphorus that plants remove from the soil.

REACTIVITY WITH PHOSPHORUS

 

1. Reaction with air

White phosphorus is very reactive. It has an appreciable vapour pressure at room temperature and inflames in dry air at about 320 K or at even lower temperatures if finely divided. In air at room temperature it emits a faint green light called phosphorescence; the reaction occurring is a complex oxidation process, but this happens only at certain partial pressures of oxygen. It is necessary, therefore, to store white phosphorus under water, unlike the less reactive red and black allotropes which do not react with air at room temperature. Both red and black phosphorus burn to form oxides when heated in air, the red form igniting at temperatures exceeding 600 K, the actual temperature depending on purity. Black phosphorus does not ignite until even higher temperatures.

 

2. Reaction with acids

Hydrochloric and dilute sulphuric acids have no appreciable action at room temperature on the pure Group V elements. Concentrated sulphuric acid and nitric acid—powerful oxidizing agents—attack all the elements except nitrogen, particularly when the acids are warm. The products obtained reflect changes in stability of the oxidation states V and III of the Group V elements. Both white and red phosphorus dissolve in, for example, concentrated nitric acid to form phosphoric(V) acid, the reaction between hot acid and white phosphorus being particularly violent. Arsenic dissolves in concentrated nitric acid forming arsenic(V) acid, H₃AsO₄, but in dilute nitric acid and concentrated sulphuric acid the main product is the arsenic(III) acid, H₃AsO₃. The more metallic element, antimony, dissolves to form the (III) oxide Sb₄O₆ with moderately concentrated nitric acid, but the (V) oxide Sb₂O₅ (structure unknown) with the more concentrated acid. Bismuth, however, forms the salt bismuth(III) nitrate Bi(NO₃)₃ 5H₂O.

3As + 5HNO₃ + 2H₂O = 3H₃AsO₄ + 5NO

3Sb + 5HNO₃ = 3HSbO₃ + 5NO + H₂O

Bi + 4HNO₃ = Bi(NO₃)₃ + NO + 2H₂O

 

3. Reaction with alkalis

The change from non-metallic to metallic properties of the Group V elements as the atomic mass of the element increases is shown in their reactions with alkalis. The head element nitrogen does not react. White phosphorus, however, reacts when warmed with a concentrated solution of a strong alkali to form phosphine, a reaction which can be regarded as a disproportionation reaction of phosphorus:

P₄ + 3KOH + 3H₂O → 3KH₂PO₂ + PH₃↑

potassium phosphine

phosphinate

hypophosphite

The phosphine produced is impure and contains small quantities of diphosphane, P₂H₄.

Arsenic, unlike phosphorus, is only slightly attacked by boiling sodium hydroxide; more rapid attack takes place with the fused alkali; an arsenate(III) is obtained in both cases,

As₄ + 12OH ^– → 4AsO₃ ^3- + 6H₂↑ cf. aluminiu.

Arsine is not formed in this reaction. Antimony and bismuth do not react with sodium hydroxide.

4. Reaction with halogens

Nitrogen does form a number of binary compounds with the halogens but none of these can be prepared by the direct combination of the elements and they are dealt with below. The other Group V elements all form halides by direct combination.

White and red phosphorus combine directly with chlorine, bromine and iodine, the red allotrope reacting in each case at a slightly higher temperature. The reactions are very vigorous and white phosphorus is spontaneously inflammable in chlorine at room temperature. Both chlorine and bromine first form a trihalide:

P₄ + 6X₂ → 4PX₃   (X = Cl or Br)

2P + 3Cl₂ ® 2PCl₃

PCl₃ + 2H₂O = H₂[PO₂H] + 3HCl

PCl₅ + HOH Û POCl₃ + 2HCl

P₂O₅ + 3PCl₅ ® 5POCl₃

 

but this is converted to a pentahalide by excess of the halogen. No pentaiodide is known.

ARSENIC, ANTIMONY AND BISMUTH

None of the common allotropic forms of these metals is affected by air unless they are heated, when all burn to the (III) oxide. A complete set of trihalides for arsenic, antimony and bismuth can be prepared by the direct combination of the elements although other methods of preparation can sometimes be used. The vigour of the direct combination reaction for a given metal decreases from fluorine to iodine (except in the case of bismuth which does not react readily with fluorine) and for a given halogen, from arsenic to bismuth. In addition to the trihalides, arsenic and antimony form pentafluorides and antimony a pentachloride; it is rather odd that arsenic pentachloride has not yet been prepared.

History

The name bismuth is from ca. 1660s, and is of uncertain etymology. It is one of the first 10 metals to have been discovered. Bismuth appears in the 1660s, from obsolete German Bismuth, Wismut, Wissmuth (early 16th century); perhaps related to Old High German hwiz (“white”). The New Latin bisemutum (due to Georgius Agricola, who Latinized many German mining and technical words) is from the German Wismuth, perhaps from weiße Masse, “white mass.” The element was confused in early times with tin and lead because of its resemblance to those elements. Bismuth has been known since ancient times, so no one person is credited with its discovery. Agricola, in De Natura Fossilium (ca. 1546) states that bismuth is a distinct metal in a family of metals including tin and lead. This was based on observation of the metals and their physical properties. Miners in the age of alchemy also gave bismuth the name tectum argenti, or “silver being made,” in the sense of silver still in the process of being formed within the Earth. Beginning with Johann Heinrich Pott in 1738, Carl Wilhelm Scheele and Torbern Olof Bergman the distinctness of lead and bismuth became clear and Claude François Geoffroy demonstrated in 1753 that this metal is distinct from lead and tin. Bismuth was also known to the Incas and used (along with the usual copper and tin) in a special bronze alloy for knives.

The word arsenic was borrowed from the Syriac word ܠܐ ܙܐܦܢܝܐ (al) zarniqa and the Persian word زرنيخ Zarnikh, meaning “yellow orpiment”, into Greek as arsenikon (Αρσενικόν). It is also related to the similar Greek word arsenikos (Αρσενικός), meaning “male”, “masculine” or “potent”. The word was adopted in Latin arsenicum and Old French arsenic, from which the English word arsenic is derived. Arsenic sulfides (orpiment, realgar) and oxides have been known and used since ancient times. Zosimos (circa 300 AD) describes roasting sandarach (realgar) to obtain cloud of arsenic (arsenious oxide), which he then reduces to metallic arsenic. As the symptoms of arsenic poisoning were somewhat ill-defined, it was frequently used for murder until the advent of the Marsh test, a sensitive chemical test for its presence. (Another less sensitive but more general test is the Reinsch test.) Owing to its use by the ruling class to murder one another and its potency and discreetness, arsenic has been called the Poison of Kings and the King of Poisons.

During the Bronze Age, arsenic was often included in bronze, which made the alloy harder (so-called “arsenical bronze”). Albertus Magnus (Albert the Great, 1193–1280) is believed to have been the first to isolate the element from a compound in 1250, by heating soap together with arsenic trisulfide. In 1649, Johann Schröder published two ways of preparing arsenic. Crystals of elemental (native) arsenic are found iature, although rare.

Cadet’s fuming liquid (impure cacodyl), often claimed as the first synthetic organometallic compound, was synthesized in 1760 by Louis Claude Cadet de Gassicourt by the reaction of potassium acetate with arsenic trioxide.

In the Victorian era, “arsenic” (“white arsenic” or arsenic trioxide) was mixed with vinegar and chalk and eaten by women to improve the complexion of their faces, making their skin paler to show they did not work in the fields. Arsenic was also rubbed into the faces and arms of women to “improve their complexion”. The accidental use of arsenic in the adulteration of foodstuffs led to the Bradford sweet poisoning in 1858, which resulted in approximately 20 deaths.

Two pigments based on arsenic have been widely used since their discovery – Paris Green and Scheele’s Green. After arsenic’s toxicity became widely known, they were less often used as pigments, so these compounds were more often used as insecticides. In the 1860s an arsenic by-product of dye production, London Purple – a solid consisting of a mixture of arsenic trioxide, aniline, lime and ferrous oxide, which is insoluble in water and very toxic by inhalation and ingestion– was widely used, but Paris Green, another arsenic based dye, was later substituted for it. With better understanding of the toxicology mechanism, two other compounds were used starting in the 1890s. Arsenite of lime and arsenate of lead were used widely as insecticides until the discovery of DDT in 1942.

Antimony(III) sulfide, Sb₂S₃, was recognized in predynastic Egypt as an eye cosmetic (kohl) as early as about 3100 BC, when the cosmetic palette was invented.

An artifact, said to be part of a vase, made of antimony dating to about 3000 BC was found at Telloh, Chaldea (part of present-day Iraq), and a copper object plated with antimony dating between 2500 BC and 2200 BC has been found in Egypt. Austen, at a lecture by Herbert Gladstone in 1892 commented that “we only know of antimony at the present day as a highly brittle and crystalline metal, which could hardly be fashioned into a useful vase, and therefore this remarkable ‘find’ (artifact mentioned above) must represent the lost art of rendering antimony malleable.”

Moorey was unconvinced the artifact was indeed a vase, mentioning that Selimkhanov, after his analysis of the Tello object (published in 1975), “attempted to relate the metal to Transcaucasiaatural antimony” (i.e. native metal) and that “the antimony objects from Transcaucasia are all small personal ornaments.” This weakens the evidence for a lost art “of rendering antimony malleable.”

The first European description of a procedure for isolating antimony is in the book De la pirotechnia of 1540 by Vannoccio Biringuccio; this predates the more famous 1556 book by Agricola, De re metallica. In this context Agricola has been often incorrectly credited with the discovery of metallic antimony. The book Currus Triumphalis Antimonii (The Triumphal Chariot of Antimony), describing the preparation of metallic antimony, was published in Germany in 1604. It was purported to have been written by a Benedictine monk, writing under the name Basilius Valentinus, in the 15th century; if it were authentic, which it is not, it would predate Biringuccio.

Pure antimony was well known to Jābir ibn Hayyān in the 8th century. There is an ongoing controversy, with translator Marcellin Berthelot stating antimony was never found in Jābir’s books, but others claiming. that Berthelot translated only some of the less important books, while the more interesting ones (some of which might describe antimony) are not yet translated, and their content is completely unknown

The first natural occurrence of pure antimony in the Earth’s crust was described by the Swedish scientist and local mine district engineer Anton von Swab in 1783; the type-sample was collected from the Sala Silver Mine in the Bergslagen mining district of Sala, Västmanland, Sweden.

Etymology of Antimony

The ancient words for antimony mostly have, as their chief meaning, kohl, the sulfide of antimony. Pliny the Elder, however, distinguishes between male and female forms of antimony; the male form is probably the sulfide, while the female form, which is superior, heavier, and less friable, has been suspected to be native metallic antimony.

The Egyptians called antimony mśdmt; in hieroglyphs, the vowels are uncertain, but there is an Arabic tradition that the word is ميسديميت mesdemet.

Pliny also gives the names stimi [sic], larbaris, alabaster, and the “very common” platyophthalmos, “wide-eye” (from the effect of the cosmetic). Later Latin authors adapted the word to Latin as stibium. The Arabic word for the substance, as opposed to the cosmetic, can appear as إثمد ithmid, athmoud, othmod, or uthmod. Littré suggests the first form, which is the earliest, derives from stimmida, an accusative for stimmi.

The use of Sb as the standard chemical symbol for antimony is due to Jöns Jakob Berzelius, who used this abbreviation of the name stibium. The medieval Latin form, from which the modern languages and late Byzantine Greek take their names for antimony, is antimonium. The origin of this is uncertain; all suggestions have some difficulty either of form or interpretation. The popular etymology, from ἀντίμοναχός anti-monachos or French antimoine, still has adherents; this would mean “monk-killer”, and is explained by many early alchemists being monks, and antimony being poisonous.

Another popular etymology is the hypothetical Greek word ἀντίμόνος antimonos, “against aloneness”, explained as “not found as metal”, or “not found unalloyed”. Lippmann conjectured a hypothetical Greek word ανθήμόνιον anthemonion, which would mean “floret”, and cites several examples of related Greek words (but not that one) which describe chemical or biological efflorescence.

The early uses of antimonium include the translations, in 1050–1100, by Constantine the African of Arabic medical treatises. Several authorities believe antimonium is a scribal corruption of some Arabic form; Meyerhof derives it from ithmid; other possibilities include athimar, the Arabic name of the metalloid, and a hypothetical as-stimmi, derived from or parallel to the Greek.

 

COMPOUNDS OF GROUP V ELEMENTS

1. HYDRIDES

All Group V elements form covalent hydrides MH₃. Some physical data for these hydrides are given below in Table 2. The abnormal values of the melting and boiling points of ammonia are explained by hydrogen bonding. The thermal stabilities of the hydrides decrease rapidly from ammonia to bismuthine as indicated by the mean thermochemical bond energies of the M—H bond, and both stibine, SbH₃, and bismuthine, BiH₃, are very unstable. All the

 

Table 2.

PROPERTIES OF GROUP V HYDRIDES

 

Group V hydrides are reducing agents, the reducing power increasing from NH₃ to BiH₃, as thermal stability decreases. These stability changes are in accordance with the change from a non-metal to a weak metal for the Group V elements nitrogen to bismuth. Nitrogen, phosphorus and arsenic form more than one hydride. Nitrogen forms several but of these only ammonia, NH₃, hydrazine, N₂H₄ and hydrogen azide N₃H (and the ammonia derivative hydroxylamine) will be considered. Phosphorus and arsenic form the hydrides diphosphane P₂H₄ and diarsane As₂H₄ respectively, but both of these hydrides are very unstable.

Hydrides of phosphorus

PHOSPHINE

Phosphine can be prepared by the reaction of a strong alkali with white phosphorus; potassium, sodium and barium hydroxides may be used:

P₄ + 3KOH + 3H₂O → 3KH₂PO₂ + PH₃↑

potassium

phosphinate

hypophosphite

This reaction gives an impure product containing hydrogen and another hydride, diphosphane, P₂H₄. Pure phosphine can be prepared by the reduction of a solution of phosphorus trichloride in dry ether with lithium aluminium hydride:

4PC1₃ + 3LiAlH₄ → 4PH₃ + 3LiCl + 3A1C1₃

Mg₃P₂ + 6H₂O = 3Mg(OH)₂ + PH₃.

2Р₄ + Ва(ОН)₂ + Н₂О = 3Ва(РО₂Н₂)₂ + 2РН₃­

 

The reaction of potassium hydroxide solution with phosphonium iodide also gives pure phosphine:

PH₄I + KOH → KI + H₂O + PH₃

Properties

Phosphine is a colourless gas at room temperature, boiling point 183K, with an unpleasant odour; it is extremely poisonous. Like ammonia, phosphine has an essentially tetrahedral structure with one position occupied by a lone pair of electrons. Phosphorus, however, is a larger atom thaitrogen and the lone pair of electrons on the phosphorus are much less ‘concentrated’ in space. Thus phosphine has a very much smaller dipole moment than ammonia. Hence phosphine is not associated (like ammonia) in the liquid state (see data in Table 2) and it is only sparingly soluble in water. Phosphine has a much lower thermal stability than ammonia and sparking decomposes it to red phosphorus and hydrogen, 2 volumes of phosphine giving 3 volumes of hydrogen. Pure phosphine ignites in air at 423 K and burns to phosphoric(V) acid :

PH₃ + 2O₂ → H₃PO₄

Applications

·        Organophosphorus chemistry

Phosphine is mainly consumed as an intermediate in organophosphorus chemistry. In an illustrative reaction, formaldehyde adds in the presence of hydrogen chloride to give tetrakis(hydroxymethyl)phosphonium chloride, which is used in textiles.

·        Microelectronics

Small amounts of phosphine are used as a dopant in the semiconductor industry, and a precursor for the deposition of compound semiconductors.

·        Fumigant

For farm use, pellets of aluminium phosphide, calcium phosphide, or zinc phosphide release phosphine upon contact with atmospheric water or rodents’ stomach acid. These pellets also contain agents to reduce the potential for ignition or explosion of the released phosphine.

Because the previously popular fumigant methyl bromide has been phased out in some countries under the Montreal Protocol, phosphine is the only widely used, cost-effective, rapidly acting fumigant that does not leave residues on the stored product. Pests developing high levels of resistance toward phosphine have become common in Asia, Australia and Brazil. High level resistance is also likely to occur in other regions, but may not have been as closely monitored.

Hydrides of arsenic and antimony

Arsine, AsH₃, and stibine, SbH₃, are formed when arsenic and antimony compounds respectively are reduced by a process in which hydrogen is evolved. They are colourless, unpleasant smelling, poisonous gases. Stibine is less stable than arsine but both decompose readily on heating to form the element and hydrogen. Both arsine and stibine are covalent compounds and they have little power to donate electrons; although the arsonium ion, AsH₄⁺, is known, this forms no stable compounds. The donor ability of arsine is enhanced when the hydrogen atoms are replaced by methyl groups.

2AsH₃ = 2As¯ + 3H₂­

Mg₃E₂ + 6HCl = 3MgCl₂ + 2H₃E

 

2. OXIDES

The principal oxides formed by Group V elements and their formal oxidation states are given below:

 

 

Oxides of phosphorus

Phosphorus forms a number of oxides, the best established being phosphorus(III) oxide, P₄O₆, and phosphorus(V) oxide, P₄O₁₀, The 4- 5 oxide is the more stable and the + 3 oxide is easily oxidised.

PHOSPHORUS(III) OXIDE, P₄O₆

Phosphorus(III) oxide is prepared by passing a slow (i.e. limited) stream of air over burning white phosphorus. A mixture of the two oxides P₄O₆ and P₄O₁₀ is thereby formed; the (V) oxide can be condensed out of the emerging gas stream as a solid by passing through a U tube heated in a water bath to about 330 K; the more volatile (III) oxide passes on and can be condensed in a second U trap surrounded by ice. Phosphorus(III) oxide dissolves slowly in cold water to yield phosphoric(III) acid, H₃PO₃ (phosphorousacid):

P₄O₆ + 6H₂O → 4H₃PO₃

With hot water a vigorous but complex reaction occurs, the products including phosphine and phosphoric(V) acid. This disproportionation reaction can be approximately represented as:

P₄O₆ + 6H₂O → PH₃ + 3H₃PO₄

PHOSPHORUS(V) OXIDE, P₄O₁₀

This oxide was originally given the formula P₂O₅ and called “phosphorus pentoxide‘; but the vapour density and structure indicate the formula P₄O₁₀.

 

P₄O₁₀ molecule

 

It is prepared by burning phosphorus in a plentiful supply of air or oxygen:

P₄ + 5O₂ → P₄O₁₀

P₂O₅ + H₂O = 2HPO₃

It is a white, deliquescent solid, very powdery, which exhibits polymorphism; on heating, several different crystalline forms appear over definite ranges of temperature—iiltimately, the P₄O₁₀ unit in the crystal disappears and a polymerised glass is obtained, which melts to a clear liquid. The most important property of phosphorus(V) oxide is its great tendency to react with water, either free or combined. It reacts with ordinary water with great vigour, and much heat is evolved; trioxophosphoric(V) acid is formed, but the local heating may convert some of this to tetraoxophosphoric(V) acid:

P₄O₁₀ + 2H₂O → 4HPO₃

HPO₃ + H₂O → H₃PO₄

Phosphorus(V) oxide will remove water from acids to give the acid anhydride. For example, if nitric acid is distilled with it, dinitrogen pentoxide is formed:

P₄O₁₀ + 4HNO₃ → 2N₂O₅ + 4HPO₃

Phosphorus(V) oxide is an extremely effective desiccating agent, reducing the vapour pressure of water over it to a negligibly small value. However, in the presence of water vapour the line powder soon becomes covered with a layer of glassy trioxophosphoric acid, and this reduces the rate at which drying can occur. For this reason, gases are better dried by passing them through loosely-packed k pentoxide rather than merely over the surface.

Oxides of arsenic

Arsenic forms two important oxides, As₄O₆ and As₄O₁₀.

ARSENIC(lIl) OXIDE, As₄O₆

This is formed when arsenic burns in air (cf. phosphorus which gives P₄O₁₀). It can exist in two crystalline modifications ; the stable one at room temperature, which also occurs naturally as arsenolite, has an octahedral form. Solid arsenic(III) oxide is easily reduced, for example by heating with charcoal, when arsenic deposits as a black shiny solid on the cooler parts of the tube.

As₂O₃ + 2H₂O = 2H₃AsO₃,

As₂O₃ + 6NaOH = 2Na₃AsO₃ + 3H₂O,

As₄O₆ + 6H₂O → 4H₃AsO₃

Arsenic(III) acid is an extremely weak acid; in fact, the oxide is amphoteric, since the following equilibria occur :

H₃AsO₃ → AsO₃^3- + 3H⁺

TOXICOLOGY OF ARSENIC(III) OXIDE

Arsenic trioxide is readily absorbed by the digestive system: toxic effects are also well known upon inhalation or upon skin contact. Elimination is rapid at first (half-life of 1–2 days), by methylation to monomethylarsonic acid and dimethylarsonic acid, and excretion in the urine, but a certain amount (30–40% in the case of repeated exposure) is incorporated into the bones, muscles, skin, hair and nails (all tissues rich in keratin) and eliminated over a period of weeks or months.

The first symptoms of acute arsenic poisoning by ingestion are digestive problems: vomiting, abdominal pains, diarrhea often accompanied by bleeding. Sub-lethal doses can lead to convulsions, cardiovascular problems, inflammation of the liver and kidneys and abnormalities in the coagulation of the blood. These are followed by the appearance of characteristic white lines (Mees stripes) on the nails and by hair loss. Lower doses lead to liver and kidney problems and to changes in the pigmentation of the skin. Even dilute solutions of arsenic trioxide are dangerous on contact with the eyes.

Chronic arsenic poisoning is known as arsenicosis. This disorder affects workers in smelters, in populations whose drinking water contains high levels of arsenic (0.3–0.4 ppm), and in patients treated for long periods with arsenic-based pharmaceuticals. Similarly, studies on workers exposed in copper foundries in the U.S., Japan and Sweden indicate a risk of lung cancer 6–10 times higher for the most exposed workers compared with the general population. Long-term ingestion of arsenic trioxide either in drinking water or as a medical treatment can lead to skin cancer. Reproductive problems (high incidence of miscarriage, low birth weight, congenital deformations) have also been indicated in one study of women exposed to arsenic trioxide dust as employees or neighbours of a copper foundry.

In Austria, there lived the so-called “arsenic eaters of Styria”, who ingested doses far beyond the lethal dose of arsenic trioxide without any apparent harm. Arsenic is thought to enable strenuous work at high altitudes, e.g. in the Alps.

ARSENIC (V) OXIDE, As₄O₁₀

Unlike phosphorus pentoxide, this oxide cannot be made directly. Arsenic(V) acid, H₃AsO₄ (strictly, tetraoxoarsenic acid), is first prepared by oxidising arsenic(III) oxide with concentrated nitric acid or some other strong oxidising agent:

2H₃AsO₃ + 2HNO₃ → 2H₃AsO₄ + NO + NO₂ + H₂O

Oxides of antimony

Antimony forms both a + 3 and a + 5 oxide. The + 3 oxide can be prepared by the direct combination of the elements or by the action of moderately concentrated nitric acid on antimony. It is an amphoteric oxide dissolving in alkalis to give antimonates(III) (for example sodium ‘antimonite’, NaSbO₂), and in some acids to form salts, for example with concentrated hydrochloric acid the trichloride, SbCl₃, is formed.

Native massive antimony with oxidation products

 

Antimony(V) oxide can be prepared by treating antimony with concentrated nitric acid. It is an oxidising agent and when gently heated loses oxygen to form the trioxide. (The change in oxidation state stability shown by antimony should be noted since it corresponds to increasing metallic character.) Unlike the amphoteric +3 oxide, the +5 oxide is acidic and dissolves only in alkalis to give hydroxoantimonates which contain the ion [Sb(OH)₆]^~. A third oxide, Sb₂O₄, is known but contains both antimony(III) and antimony(V), Sb^III(Sb^vO₄), cf. Pb₃O₄.

Sb₂O₃ + 6NaOH = 2Na₃SbO₃ + 3H₂O

 

Oxides of bismuth

Bismuth forms both + 3 and + 5 oxides. The + 3 oxide, unlike the corresponding oxides of the other Group V elements, is insoluble in alkalis, and dissolves only in acids (when bismuth salts are formed), a clear indication of the more metallic nature of bismuth. Bismuth(V) oxide is not easy to prepare; the (III) oxide (or better a suspension of the hydroxide) must be oxidised with a strong oxidising agent such as the peroxodisulphate ion. When this is carried out, the bismuthate ion, [Bi^v(OH)₆]^–, is formed. On evaporation, the sodium salt, for example, has the formula NaBiO₃. Addition of acid to a solution of a bismuthate precipitates the (V) oxide, Bi₂O₅, but this loses oxygen rapidly and forms the trioxide. The bismuthate ion is an extremely strong oxidising agent, for example the manganese(II) ion Mn²⁺ is oxidised to manganate(VII) MnO₄^–.

Bi₂O₃ + 3H₂SO₄ = Bi₂(SO₄)₃ + 3H₂O

OXOACIDS AND THEIR SALTS

Phosphorus

Phosphorus forms a large number of oxoacids, many of which cannot be isolated but do form stable salts. In general, ionisable hydrogen is bonded to the phosphorus through an oxygen atom; hydrogen atoms attached directly to phosphorus are not ionisable.

 

THE + 3 ACIDS

Two of these are important:

HPH₂O₂ phosphinic (hypophosphorous) acid and H₂PHO₃ phosphonic (orthophosphorous) acid

Phosphinic acid is a moderately strong monobasic acid. On heating the acid and its salts they disproportionate evolving phosphine:

4H₂PO₂^~ → 2PH₃ + 2HPO₄^2-

Phosphonic acid, H₃PO₃, often called just ‘phosphorous acid’, is prepared by the hydrolysis of phosphorus trichloride; a stream of air containing phosphorus trichloride vapour is passed into ice-cold water, and crystals of the solid acid separate:

PC1₃ + 3H₂O → H₃PO₃ + 3HC1

4H₃PO₃ → PH₃ + 3H₃PO₄

THE + 5 ACIDS

The important phosphoric acids and their relation to the anhydride P₄O₁₀ are:hot

P₄O₁₀   ↔ H₂0↔  HPO₃ ↔ H₂0↔   H₄P₂O₇ ↔ H₂0↔   H₃PO₄

 

H₃PO₄ ÛH⁺ + H₂PO₄^– , К₁ = 7,1×10^-3

H₂PO₄^– Û H⁺ + HPO₄^2-, К₂ + 6,2×10^-8

HPO₄^2- Û H⁺ + PO₄^3-, К₃ = 5,0×10^-13

 

(The formulae P₄O₁₀ * H₂O are merely to illustrate the interrelationship and have no structural meaning.) Tetraoxophosphoric acid, H₃PO₄:

This is prepared in the laboratory either by dissolving phosphorus(V) oxide in water (giving trioxophosphoric acid) and then heating to give the tetraoxo-acid; or by heating violet phosphorus with 33% nitric acid, which oxidises it thus:

4P + 10HNO₃ + H₂O → 4H₃PO₄ + 5NO↑ + 5NO₂↑

Tetraoxophosphoric acid is a colourless solid, very soluble in water ; an 85 % solution is often used (“syrupy phosphoric acid”).

USES OF TETRAOXOPHOSPHORIC ACID

·        As a reagent

Pure 75–85% aqueous solutions (the most common) are clear, colourless, odourless, non-volatile, rather viscous, syrupy liquids, but still pourable. Phosphoric acid is very commonly used as an aqueous solution of 85% (w/v) phosphoric acid or H₃PO_4. Because it is a concentrated acid, an 85% solution can be corrosive, although nontoxic when diluted. Because of the high percentage of phosphoric acid in this reagent, at least some of the orthophosphoric acid is condensed into polyphosphoric acids in a temperature-dependent equilibrium, but, for the sake of labeling and simplicity, the 85% represents H₃PO₄ as if it were all orthophosphoric acid. Other percentages are possible too, even above 100%, where the phosphoric acids and water would be in an unspecified equilibrium, but the overall elemental mole content would be considered specified. When aqueous solutions of phosphoric acid and/or phosphate are dilute, they are in or will reach an equilibrium after a while where practically all the phosphoric/phosphate units are in the ortho- form.

·        Rust removal

Phosphoric acid may be used as a “rust converter”, by direct application to rusted iron, steel tools, or surfaces. The phosphoric acid converts reddish-brown iron(III) oxide, Fe2O3 (rust) to black ferric phosphate, FePO4.

“Rust converter” is sometimes a greenish liquid suitable for dipping (in the same sort of acid bath as is used for pickling metal), but it is more often formulated as a gel, commonly called “naval jelly”. It is sometimes sold under other names, such as “rust remover” or “rust killer”. As a thick gel, it may be applied to sloping, vertical, or even overhead surfaces.

After treatment, the black ferric phosphate coating can be scrubbed off, leaving a fresh metal surface. Multiple applications of phosphoric acid may be required to remove all rust. The black phosphate coating can also be left in place, where it will provide moderate further corrosion resistance (such protection is also provided by the superficially similar Parkerizing and blued electrochemical conversion coating processes).

·        Food additive

Food-grade phosphoric acid (additive E338) is used to acidify foods and beverages such as various colas, but not without controversy regarding its health effects. It provides a tangy or sour taste, and being a mass-produced chemical is available cheaply and in large quantities. The low cost and bulk availability is unlike more expensive seasonings that give comparable flavors, such as citric acid which is obtainable from citrus, but usually fermented by Aspergillus niger mold from scrap molasses, waste starch hydrolysates and phosphoric acid.

·        In medicine

Phosphoric acid is used in dentistry and orthodontics as an etching solution, to clean and roughen the surfaces of teeth where dental appliances or fillings will be placed. Phosphoric acid is also an ingredient in over-the-counter anti-nausea medications that also contain high levels of sugar (glucose and fructose). This acid is also used in many teeth whiteners to eliminate plaque that may be on the teeth before application.

BIOLOGICAL EFFECTS  OF TETRAOXOPHOSPHORIC ACID

In soft drinks

Phosphoric acid, used in many soft drinks (primarily cola), has been linked in epidemiological studies to (1)chronic kidney disease and (2)lower bone density.

(1) A study performed by the Epidemiology Branch of the US National Institute of Environmental Health Sciences, concludes that drinking 2 or more colas per day was associated with doubling the risk of chronic kidney disease.

(2) A study using dual-energy X-ray absorptiometry rather than a questionnaire about breakage, provides reasonable evidence to support the theory that drinking cola results in lower bone density. This study was published in the American Journal of Clinical Nutrition. A total of 1672 women and 1148 men were studied between 1996 and 2001. Dietary information was collected using a food frequency questionnaire that had specific questions about the number of servings of cola and other carbonated beverages and that also made a differentiation between regular, caffeine-free, and diet drinks. The paper cites significant statistical evidence to show that women who consume cola daily have lower bone density. Total phosphorus intake was not significantly higher in daily cola consumers than ionconsumers; however, the calcium-to-phosphorus ratios were lower.

On the other hand, another study suggests that insufficient intake of phosphorus leads to lower bone density. The study does not examine the effect of phosphoric acid, which binds with magnesium and calcium in the digestive tract to form salts that are not absorbed, but rather studies general phosphorus intake.

A clinical study by Heaney and Rafferty using calcium-balance methods found no impact of carbonated soft drinks containing phosphoric acid on calcium excretion. The study compared the impact of water, milk, and various soft drinks (two with caffeine and two without; two with phosphoric acid and two with citric acid) on the calcium balance of 20- to 40-year-old women who customarily consumed ~3 or more cups (680 mL) of a carbonated soft drink per day. They found that, relative to water, only milk and the two caffeine-containing soft drinks increased urinary calcium, and that the calcium loss associated with the caffeinated soft drink consumption was about equal to that previously found for caffeine alone. Phosphoric acid without caffeine had no impact on urine calcium, nor did it augment the urinary calcium loss related to caffeine. Because studies have shown that the effect of caffeine is compensated for by reduced calcium losses later in the day, Heaney and Rafferty concluded that the net effect of carbonated beverages—including those with caffeine and phosphoric acid—is negligible, and that the skeletal effects of carbonated soft drink consumption are likely due primarily to milk displacement.

Other chemicals such as caffeine (also a significant component of popular common cola drinks) were also suspected as possible contributors to low bone density, due to the known effect of caffeine on calciuria. One other study, involving 30 women over the course of a week, suggests that phosphoric acid in colas has no such effect, and postulates that caffeine has only a temporary effect, which is later reversed. The authors of this study conclude that the skeletal effects of carbonated beverage consumption are likely due primarily to milk displacement(another possible confounding factor may be an association between high soft drink consumption and sedentary lifestyle.

 

Arsenic

THE + 3 ACIDS

Arsenic(III) (arsenious) acid, H₃AsO₃.—When arsenic(III) oxide is dissolved in water the corresponding acid is formed :

As₄O₆ + 6H₂O → 4H₃AsO₃

It is an extremely weak acid but does form salts. Two kinds are known, trioxoarsenates(III), for example Na₃AsO₃, and dioxoarsenates(III), for example Cu(AsO₂)^2-

THE + 5 ACIDS

Arsenic(V) acid, H₃AsO₄ (strictly, tetraoxoarsenic(V) acid) is obtained when arsenic is oxidised with concentrated nitric acid or when arsenic(V) oxide is dissolved in water. It is a moderately strong acid which, like phosphoric(V) acid, is tribasic; arsenates(V) in general resemble phosphates(V) and are often isomorphous with them. Arsenates(V) are more powerful oxidising agents than phosphates(V) and will oxidise sulphite to sulphate, hydrogen sulphide (slowly) to sulphur and, depending on the conditions, iodide to iodine.

Antimony (Sb)

No + 3 acid is known for antimony but antimonates(III) (antimonites) formed by dissolving antimony(III) oxide in alkalis are known, for example sodium dioxoantimonate(III), NaSbO₂. The + 5 acid is known in solution and antimonates(V) can be obtained by dissolving antimony(V) oxide in alkalis. These salts contain the hexahydroxoantimonate(V) ion, [Sb(OH)₆]^~.

Bismuth (Bi)

Bismuth oxide is basic. If, however, a suspension of bismuth (III) hydroxide is oxidised with a strong oxidising agent such as the peroxodisulphate ion the hexahydroxobismuthate(V) ion [Bi^v(OH)₆]^~ is formed. Evaporation of, for example, the sodium salt, gives the trioxobismuthate(V), NaBiO₃. Bismuthates(V) are extremely powerful oxidising agents and will oxidise, for example, the manganese(II) ion to manganate(VII).

 

ALLOTROPES OF PHOSPHORUS

Solid phosphorus, arsenic and antimony exist in well known allotropic modifications. Phosphorus has three main allotropic forms, white, red and black. White phosphorus is a wax-like solid made up of tetrahedral P₄ molecules with a strained P—P—P angle of 60°; these also occur in liquid phosphorus. The reactivity of white phosphorus is attributed largely to this strained structure. The rather less reactive red allotrope can be made by heating white phosphorus at 670K for several hours; at slightly higher temperatures, – 690 K, red phosphorus sublimes, the vapour condensing to reform white phosphorus. If, however, red phosphorus is heated in a vacuum and the vapour rapidly condensed, apparently another modification, violet phosphorus, is obtained. It is probable that violet phosphorus is a polymer of high molecular weight which on heating breaks down into P₂ molecules. These on cooling normally dimerise to form P₄ molecules, i.e. white phosphorus, but in vacua link up again to give the polymerised violet allotrope. Red phosphorus may have a structure intermediate between that of violet phosphorus and white phosphorus, or it may be essentially similar to the violet species. Black phosphorus is formed when white phosphorus is heated under very high pressure (12000 atmospheres). Black phosphorus has a well-established corrugated sheet structure with each phosphorus atom bonded to three neighbours. The bonding forces between layers are weak and give rise to flaky crystals which conduct electricity, properties similar to those ol graphite. It is less reactive than either white or red phosphorus. Arsenic and antimony resemble phosphorus in having several allotropic modifications. Both have an unstable yellow allotrope. These allotropes can be obtained by rapid condensation of the vapours which presumably, like phosphorus vapour, contain As₄ and Sb₄ molecules respectively. No such yellow allotrope is known for bismuth. The ordinary form of arsenic, stable at room temperature, is a grey metallic-looking brittle solid which has some power to conduct. Under ordinary conditions antimony and bismuth are silvery white and reddish white metallic elements respectively.

Some physical properties of three forms of Phosphorus are given below

Property	White phosphorus	Red phosphorus	Black Phosphorous
Colour	White, but turns yellow on exposure	Dark red	Black
Colour	White, but turns yellow on exposure	Dark red	Black
State	Waxy solid, can be cut with knife	Brittle powder	Crystalline with greasy touch
Smell	Garlic smell	Odorless
Density	1.84	2.1	2.69
Ignition temperature	307 K	533 K	673 K
Melting point	317 K	Does not melt	860 K

 

Waxy white (yellow cut), red (granules centre left, chunk centre right), and violet phosphorus

Phosphorus exists as several forms (allotropes) that exhibit strikingly different properties. The two most common allotropes are white phosphorus and red phosphorus. Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight. Black phosphorus is obtained by heating white phosphorus under high pressures (about 12,000 standard atmospheres or 1.2 gigapascals). In appearance, properties, and structure, it resembles graphite, being black and flaky, a conductor of electricity, and has puckered sheets of linked atoms. Another allotrope is diphosphorus; it contains a phosphorus dimer as a structural unit and is highly reactive.

White phosphorus

White phosphorus exposed to air glows in the darkness

White phosphorus and related molecular forms

The most important form of elemental phosphorus from the perspective of applications and the chemical literature is white phosphorus. It consists of tetrahedral P₄ molecules, in which each atom is bound to the other three atoms by a single bond. This P₄ tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C when it starts decomposing to P₂ molecules. Solid white exists in two forms. At low-temperatures, the β form is stable. At high-temperatures the α form is predominant. These forms differ in terms of the relative orientations of the constituent P₄ tetrahedra.

White phosphorus is the least stable, the most reactive, the most volatile, the least dense, and the most toxic of the allotropes. White phosphorus gradually changes to red phosphorus. This transformation is accelerated by light and heat, and samples of white phosphorus almost always contain some red phosphorus and accordingly appear yellow. For this reason it is also called yellow phosphorus. It glows in the dark (when exposed to oxygen) with a very faint tinge of green and blue, is highly flammable and pyrophoric (self-igniting) upon contact with air and is toxic (causing severe liver damage on ingestion). Owing to its pyrophoricity, white phosphorus is used as an additive iapalm. The odour of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white “(di)phosphorus pentoxide”, which consists of P₄O₁₀ tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.

Thermolysis (cracking) of P₄ at 1100 kelvin) gives diphosphorus, P₂. This species is not stable as a solid or liquid. The dimeric unit contains a triple bond and is analogous to N₂. It can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents. At still higher temperatures, P₂ dissociates into atomic P.

Although the term phosphorescence is derived from phosphorus, the reaction that gives phosphorus its glow is properly called chemiluminescence (glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it).

Crystal structure of red phosphorus

Red phosphorus

Red phosphorus is polymeric in structure. It can be viewed as a derivative of P₄ wherein one P-P bond is broken, and one additional bond is formed with the neighbouring tetrahedron resulting in a chain-like structure. Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight. Phosphorus after this treatment is amorphous. Upon further heating, this material crystallises. In this sense, red phosphorus is not an allotrope, but rather an intermediate phase between the white and violet phosphorus, and most of its properties have a range of values. For example, freshly prepared, bright red phosphorus is highly reactive and ignites at about 300 °C, though it is still more stable than white phosphorus, which ignites at about 30 °C. After prolonged heating or storage, the color darkens (see infobox images); the resulting product is more stable and does not spontaneously ignite in air.

Violet phosphorus

Violet phosphorus is a form of phosphorus that can be produced by day-long annealing of red phosphorus above 550 °C. In 1865, Hittorf discovered that when phosphorus was recrystallized from molten lead, a red/purple form is obtained. Therefore this form is sometimes known as “Hittorf’s phosphorus” (or violet or α-metallic phosphorus).

Crystal structure of black phosphorus

Black phosphorus

Black phosphorus is the least reactive allotrope and the thermodynamically stable form below 550 °C. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite. High pressures are usually required to produce black phosphorus, but it can also be produced at ambient conditions using metal salts as catalysts.

Isotopes

Twenty-three isotopes of phosphorus are known, including all possibilities from ²⁴P up to ⁴⁶P. Only ³¹P is stable and is therefore present at 100% abundance. The half-integer nuclear spin and high abundance of ³¹P make phosphorus-31 NMR spectroscopy a very useful analytical tool in studies of phosphorus-containing samples.

Two radioactive isotopes of phosphorus have half-lives that make them useful for scientific experiments. ³²P has a half-life of 14.262 days and ³³P has a half-life of 25.34 days. Biomolecules can be “tagged” with a radioisotope to allow for the study of very dilute samples.

Radioactive isotopes of phosphorus include

·  ³²P, a beta-emitter (1.71 MeV) with a half-life of 14.3 days, which is used routinely in life-science laboratories, primarily to produce radiolabeled DNA and RNA probes, e.g. for use in Northern blots or Southern blots. Because the high energy beta particles produced penetrate skin and corneas, and because any ³²P ingested, inhaled, or absorbed is readily incorporated into bone and nucleic acids, Occupational Safety and Health Administration in the United States, and similar institutions in other developed countries require that a lab coat, disposable gloves and safety glasses or goggles be worn when working with ³²P, and that working directly over an open container be avoided in order to protect the eyes. Monitoring personal, clothing, and surface contamination is also required. In addition, due to the high energy of the beta particles, shielding this radiation with the normally used dense materials (e.g. lead), gives rise to secondary emission of X-rays via Bremsstrahlung (braking radiation). Therefore shielding must be accomplished with low density materials, e.g. Plexiglas (Lucite), other plastics, water, or (when transparency is not required), even wood.

·  ³³P, a beta-emitter (0.25 MeV) with a half-life of 25.4 days. It is used in life-science laboratories in applications in which lower energy beta emissions are advantageous such as DNA sequencing.

The majority of phosphorus-containing compounds are produced for use as fertilisers. For this purpose, phosphate-containing minerals are converted to phosphoric acid. Two distinct routes are employed, the main one being treatment of phosphate minerals with sulfuric acid. The other process utilises white phosphorus, which may be produced by reaction and distillation from very low grade phosphate sources. The white phosphorus is then oxidised to phosphoric acid and subsequently neutralised with base to give phosphate salts. Phosphoric acid obtained via white phosphorus is relatively pure and is the main source of phosphates used in detergents and other non-fertiliser applications.

Elemental phosphorus

Presently, about 1,000,000 short tons (910,000 t) of elemental phosphorus is produced annually. Calcium phosphate (phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200–1,500 °C with sand, which is mostly SiO₂, and coke (impure carbon) to produce vaporized P₄. The product is subsequently condensed into a white powder under water to prevent oxidation by air. Even under water, white phosphorus is slowly converted to the more stable red phosphorus allotrope. The chemical equation for this process when starting with fluoroapatite, a common phosphate mineral, is:

4 Ca₅(PO₄)₃F + 18 SiO₂ + 30 C → 3 P₄ + 30 CO + 18 CaSiO₃ + 2 CaF₂

Side products from this production include ferrophosphorus, a crude form of Fe₂P, resulting from iron impurities in the mineral precursors. The silicate slag is a useful construction material. The fluoride is sometimes recovered for use in water fluoridation. More problematic is a “mud” containing significant amounts of white phosphorus. Production of white phosphorus is conducted in large facilities in part because it is energy intensive. The white phosphorus is transported in molten form. Some major accidents have occurred during transportation, train derailments at Brownston, Nebraska and Miamisburg, Ohio led to large fires. The worst incident in recent times was an environmental one in 1968 when the sea became contaminated due to spillages and/or inadequately treated sewage from a white phosphorus plant at Placentia Bay, Newfoundland.

Another process by which elemental phosphorus is extracted includes applying at high temperatures (1500 °C):

· 2 Ca₃(PO₄)₂ + 6 SiO₂ + 10 C → 6CaSiO₃ + 10 CO + P₄

 

Oxoacids of phosphorus

Phosphorous oxoacids are extensive, often commercially important, and sometimes structurally complicated. They all have acidic protons bound to oxygen atoms and some have nonacidic protons that are bonded directly to phosphorus. Although many oxoacids of phosphorus are formed, only nine are important, and three of them, hypophosphorous acid, phosphorous acid, and phosphoric acid, are particularly important ones.

Phosphorus(V) compounds

Oxides

The most prevalent compounds of phosphorus are derivatives of phosphate (PO₄^3-), a tetrahedral anion. Phosphate is the conjugate base of phosphoric acid, which is produced on a massive scale for use in fertilisers. Being triprotic, phosphoric acid converts stepwise to three conjugate bases:

H₃PO₄ + H₂O H₃O⁺ + H₂PO₄⁻       K_a1= 7.25×10⁻³

H₂PO₄⁻ + H₂O H₃O⁺ + HPO₄²⁻       K_a2= 6.31×10⁻⁸

HPO₄²⁻ + H₂O H₃O⁺ +  PO₄³⁻        K_a3= 3.98×10⁻¹³

Phosphate exhibits the tendency to form chains and rings with P-O-P bonds. Many polyphosphates are known, including ATP. Polyphosphates arise by dehydration of hydrogen phosphates such as HPO₄^2- and H₂PO₄^–. For example, the industrially important trisodium triphosphate (also known as sodium tripolyphosphate, STPP) is produced industrially on a megatonne scale via this condensation reaction:

2 Na₂[(HO)PO₃] + Na[(HO)₂PO₂] → Na₅[O₃P-O-P(O)₂-O-PO₃] + 2 H₂O

Phosphorus pentoxide (P₄O₁₀) is the acid anhydride of phosphoric acid, but several intermediates are known between the two. This waxy white solid reacts vigorously with water.

The tetrahedral structure of P₄O₁₀ and P₄S₁₀

With metal cations, phosphate forms a variety of salts. These solids are polymeric, featuring P-O-M linkages. When the metal cation has a charge of 2+ or 3+, the salts are generally insoluble, hence they exist as common minerals. Many phosphate salts are derived from hydrogen phosphate (HPO₄^2-).

PCl₅ and PF₅ are common compounds. Both are volatile and pale or colourless. The other two halides, PBr₅ and PI₅PI₅ are unstable. The pentachloride and pentafluoride adopt trigonal bipyramid molecular geometry and are Lewis acids. With fluoride, PF₅ forms PF₆^–, an anion that is isoelectronic with SF₆. The most important oxyhalide is phosphorus oxychloride (POCl₃), which is tetrahedral.

Before extensive computer calculations were feasible, it was proposed that bonding in phosphorus(V) compounds involved d orbitals. It is now accepted that the bonding can be better explained by molecular orbital theory and involves only s- and p-orbitals on phosphorus.

Nitrides

Compounds of the formula (PNCl₂)_n exist mainly as rings such as the trimer hexachlorophosphazene. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride: PCl₅ + NH₄Cl → 1/n (NPCl₂)_n + 4 HCl The chloride groups can be replaced by alkoxide (RO^–) to give rise to a family of polymers with potentially useful properties.

Sulfides

Main article: phosphorus sulfide

Phosphorus forms a wide range of sulfides, where phosphorus can be P(V), P(III) or other oxidation states. Most famous is the three-fold symmetric P₄S₃ used in strike-anywhere matches. P₄S₁₀ and P₄O₁₀ have analogous structures.

Phosphorus(III) compounds

Phosphine (PH₃) and its organic derivatives (PR₃) are structural analogues with ammonia (NH₃) but the bond angles at phosphorus are closer to 90° for phosphine and its organic derivatives. It is an ill-smelling, toxic compound. Phosphine is produced by hydrolysis of calcium phosphide, Ca₃P₂. Unlike ammonia, phosphine is oxidised by air. Phosphine is also far less basic than ammonia.

All four symmetrical trihalides are well known: gaseous PF₃, the yellowish liquids PCl₃ and PBr₃, and the solid PI₃. These materials are moisture sensitive, hydrolysing to give phosphorus acid. The trichloride, a common reagent, is produced by chlorination of white phosphorus:

P₄ + 6 Cl₂ → 4 PCl₃

The trifluoride is produced by from the trichloride by halide exchange. PF₃ is toxic because it binds to haemoglobin.

Phosphorus(III) oxide, P₄O₆ (also called tetraphosphorus hexoxide) is the anhydride of P(OH)₃, the minor tautomer of phosphorous acid. The structure of P₄O₆ is like that of P₄O₁₀ less the terminal oxide groups.

Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

Organophosphorus compounds

Main article: organophosphorus compounds

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl₃ serves as a source of P³⁺ in routes to organophosphorus(III) compounds. For example it is the precursor to triphenylphosphine:

PCl₃ + 6 Na + 3 C₆H₅Cl → P(C₆H₅)₃ + 6 NaCl

Treatment of phosphorus trihalides with alcohols and phenols gives phosphites, e.g. triphenylphosphite:

PCl₃ + 3 C₆H₅OH → P(OC₆H₅)₃ + 3 HCl

Similar reactions occur for phosphorus oxychloride, affording triphenylphosphate:

OPCl₃ + 3 C₆H₅OH → OP(OC₆H₅)₃ + 3 HCl

Phosphorus(I) and phosphorus(II) compounds

These compounds generally feature P-P bonds. Examples include catenated derivatives of phosphine and organophosphines. The highly flammable gas diphosphine (P₂H₄) is the first of a series of derivatives of this type. Diphosphine is an analogue of hydrazine. Compounds containing P=P double bonds have also been observed, although they are rare.

A stable diphosphene, a derivative of phosphorus(I).

Phosphides

The phosphide ion is P^3-. Phosphides arise by reaction of metals with red phosphorus. Salts of P^3- do not exist in solution and these derivatives are refractory, reflecting their high lattice energy. Illustrated by the behaviour calcium phosphide, many metal phosphides hydrolyse in water with release of phosphine:

Ca₃P₂ + 6 H₂O → 2 PH₃ + 3 Ca(OH)₂

Schreibersite is a naturally occurring phosphide found in meteorites. Many polyphosphides are known such as derivatives of OsP₂. These can be structurally complex ranging from Na₃P₇ and derivatives of P₂₆^4-. Often these species adopt cage-like structures that resemble fragments of violet phosphorus.

Spelling and etymology

The name Phosphorus in Ancient Greece was the name for the planet Venus and is derived from the Greek words (φως = light, φέρω = carry), which roughly translates as light-bringer or light carrier. (In Greek mythology and tradition, Augerinus (Αυγερινός = morning star, in use until today), Hesperus or Hesperinus (΄Εσπερος or Εσπερινός or Αποσπερίτης = evening star, in use until today) and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity) are close homologues, and also associated with Phosphorus-the-planet).

According to the Oxford English Dictionary, the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P³⁺ valence: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (e.g., phosphorous acid) and P⁵⁺ valence phosphoric compounds (e.g., phosphoric acids and phosphates).

 

Applications

Fertiliser

The dominant application of phosphorus is in fertilisers, which provides phosphate as required for all life and is often a limiting nutrient for crops. Phosphorus, being an essential plant nutrient, finds its major use as a constituent of fertilizers for agriculture and farm production in the form of concentrated phosphoric acids, which can consist of 70% to 75% P₂O₅. Global demand for fertilisers led to large increase in phosphate (PO₄^3–) production in the second half of the 20th century. Due to the essential nature of phosphorus to living organisms, the low solubility of natural phosphorus-containing compounds, and the slow natural cycle of phosphorus, the agricultural industry relies on fertilisers that contain phosphate. A major form of these fertilisers is superphosphate of lime, a mixture of two salts, calcium dihydrogen phosphate Ca(H₂PO₄)₂ and calcium sulfate dihydrate CaSO₄·2H₂O, produced by the reaction of sulfuric acid and water with calcium phosphate.

Organophosphorus compounds

White phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and two phosphorus sulfides, phosphorus pentasulfide, and phosphorus sesquisulfide. Organophosphorus compounds have many applications, including in plasticizers, flame retardants, pesticides, extraction agents, and water treatment. in particular the herbicide glyphosate sold under the brand name Roundup.

Metallurgical aspects

Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products. Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce phosphorus-containing copper (CuOFP) alloys with a higher hydrogen embrittlement resistance thaormal copper.

Matches

Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction

 

Phosphorus-contained matches were first produced in 1830s and contained a mixture of white phosphorus, an oxygen-releasing compound (potassium chlorate, lead dioxide or some nitrate) and a binder in their heads. They were rather sensitive to storage conditions, toxic and unsafe, as they could be lit by striking on any rough surface. Therefore, their production was gradually banned between 1872 and 1925 in different countries. The international Berne Convention, adopted in 1906, prohibited the use of white phosphorus in matches. As a consequence, the ‘strike-anywhere’ matches were gradually replaced by ‘safety matches’ where white phosphorus was substituted by phosphorus sesquisulfide (P₄S₃), sulfur or antimony sulfide. Such matches are hard to ignite on an arbitrary surface and require a special strip. The strip contains red phosphorus which heats up upon striking, reacts with the oxygen-releasing compound in the head and ignites the flammable material of the head.

Water softening

Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others. It is useful for softening water to enhance the performance of the detergents and to prevent pipe/boiler tube corrosion.

Niche applications

· Phosphates are utilized in the making of special glasses that are used for sodium lamps.

· Bone-ash, calcium phosphate, is used in the production of fine china.

· Phosphoric acid made from elemental phosphorus is used in food applications such as some soda beverages. The acid is also a starting point to make food grade phosphates. These include mono-calcium phosphate that is employed in baking powder and sodium tripolyphosphate and other sodium phosphates. Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste.

· White phosphorus, called “WP” (slang term “Willie Peter”) is used in military applications as incendiary bombs, for smoke-screening as smoke pots and smoke bombs, and in tracer ammunition. It is also a part of an obsolete M34 White Phosphorus US hand grenade. This multipurpose grenade was mostly used for signalling, smoke screens and inflammation; it could also cause severe burns and had a psychological impact on the enemy. Military uses of white phosphorus are constrained by international law

· In trace amounts, phosphorus is used as a dopant for n-type semiconductors.

· ³²P and ³³P are used as radioactive tracers in biochemical laboratories.

· Phosphate is a strong complexing agent for the hexavalent uranyl (UO₂²⁺) species and this is the reason why apatite and other natural phosphates can often be very rich in uranium.

· Tributylphosphate is an organophosphate soluble in kerosene and used to extract uranium in the Purex process applied in the reprocessing of spent nuclear fuel.

Biological role

Inorganic phosphorus in the form of the phosphate PO³⁻₄ is required for all known forms of life, playing a major role in biological molecules such as DNA and RNA where it forms part of the structural framework of these molecules. Living cells also use phosphate to transport cellular energy in the form of adenosine triphosphate (ATP). Nearly every cellular process that uses energy obtains it in the form of ATP. ATP is also important for phosphorylation, a key regulatory event in cells. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts assist in stiffening bones.

Living cells are defined by a membrane that separates it from its surroundings. Biological membranes are made from a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol, such that two of the glycerol hydroxyl (OH) protons have been replaced with fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.

An average adult human contains about 0.7 kg of phosphorus, about 85–90% of which is present in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids (~1%). The phosphorus content increases from about 0.5 weight% in infancy to 0.65–1.1 weight% in adults. Average phosphorus concentration in the blood is about 0.4 g/L, about 70% of that is organic and 30% inorganic phosphates. A well-fed adult in the industrialized world consumes and excretes about 1–3 grams of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as nucleic acids and phospholipids; and excretion almost exclusively in the form of phosphate ions such as H₂PO−₄ and HPO²⁻₄. Only about 0.1% of body phosphate circulates in the blood, and this amount reflects the amount of phosphate available to soft tissue cells.

Bone and teeth enamel

The main component of bone is hydroxyapatite as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. Water fluoridation enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material called fluoroapatite:

Ca₅(PO₄)₃OH + F⁻ → Ca₅(PO₄)₃F + OH⁻

Phosphorus deficiency

In medicine, low-phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as re-feeding after malnutrition) or pass too much of it into the urine. All are characterized by hypophosphatemia, which is a condition of low levels of soluble phosphate levels in the blood serum, and therefore inside cells. Symptoms of hypophosphatemia include muscle and neurological dysfunction, and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhoea and calcification (hardening) of organs and soft tissue, and can interfere with the body’s ability to use iron, calcium, magnesium, and zinc.

Phosphorus is an essential macromineral for plants, which is studied extensively in edaphology in order to understand plant uptake from soil systems. In ecological terms, phosphorus is often a limiting factor in many environments; i.e. the availability of phosphorus governs the rate of growth of many organisms. In ecosystems an excess of phosphorus can be problematic, especially in aquatic systems, resulting in eutrophication which sometimes lead to algal blooms.

Food sources

The main food sources for phosphorus are foods containing protein, although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if one holds a meal plan providing sufficient amount of protein and calcium then the amount of phosphorus is also likely sufficient.

Precautions

Organic compounds of phosphorus form a wide class of materials, many are required for life, but some are extremely toxic. Fluorophosphate esters are among the most potent neurotoxins known. A wide range of organophosphorus compounds are used for their toxicity to certain organisms as pesticides (herbicides, insecticides, fungicides, etc.) and weaponised as nerve agents. Most inorganic phosphates are relatively nontoxic and essential nutrients.

The white phosphorus allotrope presents a significant hazard because it ignites in air and produces phosphoric acid residue. Chronic white phosphorus poisoning leads to necrosis of the jaw called “phossy jaw”. Ingestion of white phosphorus may cause a medical condition known as “Smoking Stool Syndrome”.^]

Phosphorus explosion

Upon exposure to elemental phosphorus, in the past it was suggested to wash the affected area with 2% copper sulfate solution to form harmless compounds that can be washed away. According to the recent US Navy’s Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries, “Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis.”

The manual suggests instead “a bicarbonate solution to neutralize phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots. Promptly debride the burn if the patient’s condition will permit removal of bits of WP (white phosphorus) that might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns.” As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.

 

Organophosphate poisoning results from exposure to organophosphates (OPs), which cause the inhibition of acetylcholinesterase (AChE), leading to the accumulation of acetylcholine (ACh) in the body. Organophosphate poisoning most commonly results from exposure to insecticides or nerve agents. OPs are one of the most common causes of poisoning worldwide, and are frequently intentionally used in suicides in agrarian areas. There are around 1 million OP poisonings per year with several hundred thousand resulting in fatalities annually.

Organophosphates inhibit AChE, causing OP poisoning by phosphorylating the serine hydroxyl residue on AChE, which inactivates AChE. AChE is critical for nerve function, so the irreversible blockage of this enzyme, which causes acetylcholine accumulation, results in muscle overstimulation. This causes disturbances across the cholinergic synapses and can only be reactivated very slowly, if at all. Paraoxonase (PON1) is a key enzyme involved in OP pesticides and has been found to be critical in determining an organism’s sensitivity to OP exposure.

SIGNS AND SYMPTOMS

The health effects associated with organophosphate poisoning are a result of excess acetylcholine (ACh) present at different nerves and receptors in the body because acetyocholinesterase is blocked. Accumulation of ACh at motor nerves causes overstimulation of nicotinic expression at the neuromuscular junction. When this occurs symptoms such as muscle weakness, fatigue, muscle cramps, fasciculation, and paralysis can be seen. When there is an accumulation of ACh at autonomic ganglia this causes overstimulation of nicotinic expression in the sympathetic system. Symptoms associated with this are hypertension, and hypoglycemia. Overstimulation of nicotinic acetylcholine receptors in the central nervous system, due to accumulation of ACh, results in anxiety, headache, convulsions, ataxia, depression of respiration and circulation, tremor, general weakness, and potentially coma. When there is expression of muscarinic overstimulation due to excess acetylcholine at muscarinic acetylcholine receptors symptoms of visual disturbances, tightness in chest, wheezing due to bronchoconstriction, increased bronchial secretions, increased salivation, lacrimation, sweating, peristalsis, and urination can occur.

The effects of organophosphate poisoning on muscarinic receptors are recalled using the mnemonic SLUDGEM (Salivation, Lacrimation, Urination, Defecation, Gastrointestinal motility, Emesis, miosis). An additional mnemonic is MUDDLES: miosis, urination, diarrhea, diaphoresis, lacrimation, excitation, and salivation.

The onset and severity of symptoms, whether acute or chronic, depends upon the specific chemical, the route of exposure, the dose, and the individuals ability to degrade the compound, which the PON1 enzyme level will affect.

·        Reproductive effects

Certain reproductive effects in fertility, growth, and development for males and females have been linked specifically to OP pesticide exposure. Most of the research on reproductive effects has been conducted on farmers working with pesticides and insecticdes in rural areas. For those males exposed to OP pesticides, poor semen and sperm quality have been seen, including reduced seminal volume and percentage motility, as well as a decrease in sperm count per ejacuate. In females menstrual cycle disturbances, longer pregnancies, spontaneous abortions, stillbirths, and some developmental effects in offspring have been linked to OP pesticide exposure. Prenatal exposure has been linked to impaired fetal growth and development. The effects of OP exposure on infants and children are at this time currently being researched to come to a conclusive finding. Evidence of OP exposure in pregnant mothers are linked to several health effects in the fetus. Some of these effects include delayed mental development, Pervasive developmental disorder (PDD), morphological abnormalities in the cerebral surface. Studies observing prenatal exposure to OP pesticides used Latina women living in various agricultural communities in California as their focus. In these studies, exposure to OP pesticides were mainly measured in three ways : (1) using dialkyl phosphate (DAP) metabolites collected from urine (2) using pesticide-specific metabolites found in urine and (3) measuring cholinesterase (ChE) and butyryl cholinesterase (BChE). It was found that high exposure to OP pesticides prenatally (measured by DAP in urine) is associated with lowered IQ scores in 7-year-old children. Higher DAP levels were associated with lower scores on four cognitive areas, with the strongest association in verbal comprehension. A separate study also showed that higher prenatal chlorpyrifos (CPF) exposure was also linked to several brain anomalies. This studies used MRI and compared 20 low CPF exposure children to 20 high CPF exposure children. It was found that although overall brain size did not differ between exposure groups, there were significant areas of the brain that were enlarged in children with high exposure. These enlargements were primarily due to an increase of underlying white matter in the high exposure group. Enlargements were seen bilaterally in the superior temporal, posterior middle temporal, and inferior postcentral gyri, and superior frontal gyrus, straight gyrus, cuneus, and precuneus in the medial views of the right hemisphere.

·        Neurotoxic effects

Neurotoxic effects have also been linked to poisoning with OP pesticides causing four neurotoxic effects in humans: cholinergic syndrome, intermediate syndrome, organophosphate-induced delayed polyneuropathy (OPIDP), and chronic organophosphate-induced neuropsychiatric disorder (COPIND). These syndromes result after acute and chronic exposure to OP pesticides.

Cholinergic syndrome occurs in acute poisonings with OP pesticides and is directly related to levels of AChE activity. Symptoms include miosis, sweating, lacrimation, gastrointestinal symptoms, respiratory difficulties, dyspnea, bradycardia, cyanosis, vomiting, diarrhea, as well as other symptoms. Along with these central effects can be seen and finally seizures, convulsions, coma, respiratory failure. If the person survives the first day of poisoning personality changes can occur, aggressive events, psychotic episodes, disturbances and deficits in memory and attention, as well as other delayed effects. When death occurs, it is most commonly due to respiratory failure from the combination of central and peripheral effects, paralysis of respiratory muscles and depression of the brain respiratory center. For people afflicted with cholinergic syndrome, atropine sulfate combined with an oxime is used to combat the effects of the acute OP poisoning. Diazepam is sometimes also administered in combination with the atropine and oximes.

The intermediate syndrome (IMS) appears in the interval between the end of the cholinergic crisis and the onset of OPIDP. Symptoms associated with IMS manifest within 24–96 hours after exposure. The exact etiology, incidence, and risk factors associated with IMS are not clearly understood, but IMS is recognized as a disorder of neuromuscular junctions. IMS occurs when a person has a prolonged and severe inhibition of AChE and has been linked to specific OP pesticides such as methylparathion, dichlorvos, and parathion. Patients present with increasing weakness of facial, neck flexor and respiratory muscles.

OPIDP occurs in a small percentage of cases, roughly two weeks after exposure, where temporary paralysis occurs. This loss of function and ataxia of peripheral nerves and spinal cord is the phenomenon of OPIDP. Once the symptoms begin with shooting pains in both legs, the symptoms continue to worse for 3–6 months. In the most severe cases quadriplegia has been observed. Treatment only affects sensory nerves, not motor neurons which may permanently lose function. The aging and phosphorylation of more than 70% of functional NTE in peripheral nerves is one of the processes involved in OPIDP. Standard treatments for OP poisoning are ineffective for OPIDP.

COPIND occurs without cholinergic symptoms and is not dependent on AChE inhibition. COPIND appears with a delay and is long lasting. Symptoms associated with COPIND include cognitive deficit, mood change, autonomic dysfunction, peripheral neuropathy, and extrapyramidal symptoms. The underlying mechanisms of COPIND have not been determined, but it is hypothesized that withdrawal of OP pesticides after chronic exposure or acute exposure could be a factor.

Effects on developing animals

Evidence of exposure to OP pesticides during gestation and early postnatal period have been linked to neurodevelopmental effects in animals, specifically rats. Animals exposed in utero to chlorpyrifos exhibited decreased balance, poorer cliff avoidance, decreased locomotion, delays in maze performance, and increased gait abnormalities. Early gestation is believed to be a critical time period for the neurodevelopmental effects of pesticides. OP’s affect the cholinergic system of fetuses, so exposure to chlorpyrifos during critical periods of brain development potentially could cause cellular, synaptic, and neurobehavioral abnormalities in animals. In rats exposed to methyl parathion, studies found reduced AChE activity in all brain regions and subtle alterations in behaviors such as locomotor activity and impaired cage emergence. Organophosphates as whole have been linked to decreases in the length of limbs, head circumference, and slower rates of postnatal weight gain in mice.

 

Cause

OP pesticide exposure occurs through inhalation, ingestion and dermal contact. Because OP pesticides disintegrate quickly in air and light, they have been considered relatively safe to consumers. However, OP residues linger on fruits and vegetables. Certain OP pesticides have been banned for use on some crops, for example methyl parathion is banned from use on some crops while permitted on others. The Environmental Working Group has developed lists for concerned consumers, identifying crops with the highest pesticide residue quantities and the lowest. The “Dirty Dozen” crops are updated yearly and in 2012 included apples, celery, sweet bell peppers, peaches, strawberries, imported nectarines, grapes, spinach, lettuce, cucumbers, domestic blueberries and potatoes. Forty-five fruits and vegetables are listed by the Environmental Working Group as being regularly found with Pesticide residue associated with OPs.

Examples

·        Insecticides including malathion, parathion, diazinon, fenthion, dichlorvos, chlorpyrifos, ethion

·        Nerve gases including soman, sarin, tabun, VX

·        Ophthalmic agents: echothiophate, isoflurophate

·        Antihelmintics such as trichlorfon

·        Herbicides including tribufos [DEF], merphos are tricresyl phosphate–containing industrial chemicals.

Exposure to any one of the above listed organophosphates occurs on a daily basis through inhalation, absorption, and ingestion, most commonly of food that has been treated with an organophosphate herbicide or insecticide. Exposure to these chemicals can occur at public buildings, schools, residential areas, and in agricultural areas. The chemicals chlorpyrifos and malathion have been linked to reproductive effects, neurotoxicity, kidney/liver damage, and birth defects. Dichlorvos has also been linked to reproductive effects, neurotoxicity, and kidney/liver damage, as well as being a possible carcinogen.

 

Pathophysiology

Paraoxonase (PON1) is a key enzyme in the metabolism of organophosphates. PON1 can inactivate some OPs through hydrolysis. PON1 hydrolyzes the active metabolites in several OP insecticides such as chlorpyrifos oxon, and diazoxon, as well as, nerve agents such as soman, sarin, and VX. PON1 hydrolyzes the metabolites, not the parent compounds of insectides. The presence of PON1 polymorphisms causes there to be different enzyme levels and catalytic efficiency of this esterase, which in turn suggests that different individuals may be more susceptible to the toxic effect of OP exposure. The level of PON1 plasma hydrolytic activity provides more protection against OP pesticides. Rats injected with purified PON1 from rabbit serum were more resistant to acute cholinergic activity than the control rats. PON1 knockouts in mice are found to be more sensitive to the toxicity of pesticides, like chlorpyrifos. Animal experiments indicate that while PON1 plays a significant role in regulating the toxicity of OPs its degree of protection given depends on the compound (i.e. Chlorpyrifos oxon or diazoxon). The catalytic efficiency with which PON1 can degrade toxic OPs determines the degree of protection that PON1 can provide for organism. The higher the concentration of PON1 the better the protection provided. PON1 activity is much lower ieonates, so neonates are more sensitive to OP exposure. In 2006, reports up to a 13-fold variation was seen in PON1 levels in adults, as well as, specifically regarding sensitivity to diazoxon, a variation up to 26 and 14-fold was reported in a group of newborns and Latino mothers. This wide range in variability of enzyme levels determining a humans sensitivity to various OPs is being researched further.

 

Diagnosis

A number of measurements exist to assess exposure and early biological effects for organophosphate poisoning. Measurements of OP metabolites in both the blood and urine can be used to determine if a person has been exposed to organophosphates. Specifically in the blood, metabolites of cholinesterases, such as butyrylcholinesterase (BuChE) activity in plasma, neuropathy target esterase (NTE) in lymphocytes, and of acetylcholinesterase (AChE) activity in red blood cells. Due to both AChE and BuChE being the main targets of organophosphates, their measurement is widely used as an indication of an exposure to an OP. The main restriction on this type of diagnosis is that depending on the OP the degree to which either AChE or BuChE are inhibited differs; therefore, measure of metabolites in blood and urine do not specify for a certain OP. However, for fast initial screening, determining AChE and BuChE activity in the blood are the most widely used procedures for confirming a diagnosis of OP poisoning.

Treatment

Current antidotes for OP poisoning consist of a pretreatment with carbamates to protect AChE from inhibition by OP compounds and post-exposure treatments with anti-cholinergic drugs. Anti-cholinergic drugs work to counteract the effects of excess acetylcholine and reactivate AChE. Atropine can be used as an antidote in conjunction with pralidoxime or other pyridinium oximes (such as trimedoxime or obidoxime), though the use of “-oximes” has been found to be of no benefit, or possibly harmful, in at least two meta-analyses. Atropine is a muscarinic antagonist, and thus blocks the action of acetylcholine peripherally. These antidotes are effective at preventing lethality from OP poisoning, but current treatment lack the ability to prevent post-exposure incapacitation, performance deficits, or permanent brain damage.

Enzyme bioscavengers are being developed as a pretreatment to sequester highly toxic OPs before they can reach their physiological targets and prevent the toxic effects from occurring. Significant advances with cholinesterases (ChEs), specifically human serum BChE (HuBChE) have been made. HuBChe can offer a broad range of protection for nerve agents including soman, sarin, tabun, and VX. HuBChE also possess a very long retention time in the human circulation system and because it is from a human source it will not produce any antagonistic immunological responses. HuBChE is currently being assessed for inclusion into the protective regimen against OP nerve agent poisoning. Currently there is potential for PON1 to be used to treat sarin exposure, but recombinant PON1 variants would need to first be generated to increase its catalytic efficiency.

One other agent that is being researched is the Class III anti-arrhythmic agents. Hyperkalemia of the tissue is one of the symptoms associated with OP poisoning. While the cellular processes leading to cardiac toxicity are not well understood, the potassium current channels are believed to be involved. Class II anti-arrhythmic agents block the potassium membrane currents in cardiac cells, which makes them a candidate for become a therapeutic of OP poisoning.

History

·        Ginger Jake

A striking example of OPIDN occurred during the 1930s Prohibition Era when thousands of men in the American South and Midwest developed arm and leg weakness and pain after drinking a “medicinal” alcohol substitute. The drink, called “Ginger Jake,” contained an adulterated Jamaican ginger extract containing tri-ortho-cresyl phosphate (TOCP) which resulted in partially reversible neurologic damage. The damage resulted in the limping “Jake Leg” or “Jake Walk” which were terms frequently used in the blues music of the period. Europe and Morocco both experienced outbreaks of TOCP poisoning from contaminated abortifacients and cooking oil, respectively.

·        Gulf War Syndrome

Research has linked the neurological abnormalities found in Persian Gulf War veterans, who suffer from Gulf War syndrome, to exposure to wartime combinations of organophosphate chemical nerve agents. Before, it was believed that veterans were suffering from a psychologically based disorder or depression, most likely post-traumatic stress disorder (PTSD). Many veterans were given pyridostigmine bromide (PB) pills to protect against nerve gas agents such as sarin and soman. During the war veterans were exposed to combinations of organophosphate nerve agents, which produced symptoms associated with chronic organophosphate-induced delayed polyneuropathy (OPIDP)syndrome. Similar symptoms found in the veterans were the same symptoms reported for individuals in occupational settings who were acutely poisoned by organohosphates, such as chlorpyrifos. Studies found veterans experienced deficits in intellectual and academic abilities, simple motor skills, memory impairment, and impaired emotional function. These symptoms indicate brain damage, not a psychologically based disorder.

 

Arsenic is a chemical element with symbol As and atomic number 33. Arsenic occurs in many minerals, usually in conjunction with sulfur and metals, and also as a pure elemental crystal. It was first documented by Albertus Magnus in 1250. Arsenic is a metalloid. It can exist in various allotropes, although only the gray form has important use in industry.

The main use of metallic arsenic is for strengthening alloys of copper and especially lead (for example, in car batteries). Arsenic is a commo-type dopant in semiconductor electronic devices, and the optoelectronic compound gallium arsenide is the most common semiconductor in use after doped silicon. Arsenic and its compounds, especially the trioxide, are used in the production of pesticides, treated wood products, herbicides, and insecticides. These applications are declining, however.

Arsenic is notoriously poisonous to multicellular life, although a few species of bacteria are able to use arsenic compounds as respiratory metabolites. Arsenic contamination of groundwater is a problem that affects millions of people across the world.

Physical characteristics

Crystal structure common to Sb, AsSb and gray As

The three most common arsenic allotropes are metallic gray, yellow and black arsenic, with gray being the most common. Gray arsenic (α-As, space group R3m No. 166) adopts a double-layered structure consisting of many interlocked ruffled six-membered rings. Because of weak bonding between the layers, gray arsenic is brittle and has a relatively low Mohs hardness of 3.5. Nearest and next-nearest neighbors form a distorted octahedral complex, with the three atoms in the same double-layer being slightly closer than the three atoms in the next. This relatively close packing leads to a high density of 5.73 g/cm³. Gray arsenic is a semimetal, but becomes a semiconductor with a bandgap of 1.2–1.4 eV if amorphized. Yellow arsenic is soft and waxy, and somewhat similar to tetraphosphorus (P₄). Both have four atoms arranged in a tetrahedral structure in which each atom is bound to each of the other three atoms by a single bond. This unstable allotrope, being molecular, is the most volatile, least dense and most toxic. Solid yellow arsenic is produced by rapid cooling of arsenic vapor, As₄. It is rapidly transformed into the gray arsenic by light. The yellow form has a density of 1.97 g/cm³. Black arsenic is similar in structure to red phosphorus.

Chemistry

When heated in air, arsenic oxidizes to arsenic trioxide; the fumes from this reaction have an odor resembling garlic. This odor can be detected on striking arsenide minerals such as arsenopyrite with a hammer. Arsenic (and some arsenic compounds) sublimes upon heating at atmospheric pressure, converting directly to a gaseous form without an intervening liquid state at 887 K (614 °C). The triple point is 3.63 MPa and 1,090 K (820 °C). Arsenic makes arsenic acid with concentrated nitric acid, arsenious acid with dilute nitric acid, and arsenic trioxide with concentrated sulfuric acid.

Arsenic burning in the air

Compounds

Arsenic compounds resemble in some respects those of phosphorus, which occupies the same group (column) of the periodic table. Arsenic is less commonly observed in the pentavalent state, however. The most common oxidation states for arsenic are: −3 in the arsenides, such as alloy-like intermetallic compounds; and +3 in the arsenites, arsenates(III), and most organoarsenic compounds. Arsenic also bonds readily to itself as seen in the square As3−4 ions in the mineral skutterudite. In the +3 oxidation state, arsenic is typically pyramidal, owing to the influence of the lone pair of electrons.

Inorganic

Arsenic forms colorless, odorless, crystalline oxides As₂O₃ (“white arsenic”) and As₂O₅, which are hygroscopic and readily soluble in water to form acidic solutions. Arsenic(V) acid is a weak acid. Its salts are called arsenates, which is the basis of arsenic contamination of groundwater, a problem that affects many people. Synthetic arsenates include Paris Green (copper(II) acetoarsenite), calcium arsenate, and lead hydrogen arsenate. The latter three have been used as agricultural insecticides and poisons.

The protonation steps between the arsenate and arsenic acid are similar to those between phosphate and phosphoric acid. Unlike phosphorus acid, arsenous acid is genuinely tribasic, with the formula As(OH)₃.

A broad variety of sulfur compounds of arsenic are known. Orpiment (As₂S₃) and realgar (As₄S₄) are somewhat abundant and were formerly used as painting pigments. In As₄S₁₀, arsenic has a formal oxidation state of +2 in As₄S₄, which features As-As bonds so that the total covalency of As is still three.

The trifluoride, trichloride, tribromide, and triiodide of arsenic(III) are well known, whereas only Arsenic pentafluoride (AsF₅) is the only important pentahalide. Again reflecting the lower stability of the 5+ oxidation state, the pentachloride is stable only below −50 °C.

Organoarsenic compounds

Trimethylarsine

A large variety of organoarsenic compounds are known. Several were developed as chemical warfare agents during World War I, including vesicants such as lewisite and vomiting agents such as adamsite. Cacodylic acid, which is of historic and practical interest, arises from the methylation of arsenic trioxide, a reaction that has no analogy in phosphorus chemistry.

Alloys

Arsenic is used as the group 5 element in the III-V semiconductors gallium arsenide, indium arsenide, and aluminium arsenide. The valence electron count of GaAs is the same as a pair of Si atoms, but the band structure is completely different, which results distinct bulk properties. Other arsenic alloys include the II-IV semiconductor cadmium arsenide.

Occurrence and production

A large sample of native arsenic

Minerals with the formula MAsS and MAs₂ (M = Fe, Ni, Co) are the dominant commercial sources of arsenic, together with realgar (an arsenic sulfide mineral) and native arsenic. An illustrative mineral is arsenopyrite (FeAsS), which is structurally related to iron pyrite. Many minor As-containing minerals are known. Arsenic also occurs in various organic forms in the environment. Inorganic arsenic and its compounds, upon entering the food chain, are progressively metabolized to a less toxic form of arsenic through a process of methylation.

Other naturally occurring pathways of exposure include volcanic ash, weathering of arsenic-containing minerals and ores, and dissolved in groundwater. It is also found in food, water, soil, and air. Arsenic is absorbed by all plants, but is more concentrated in leafy vegetables, rice, apple and grape juice, and seafood. An additional route of exposure is through inhalation.

On roasting in air of arsenopyrite, arsenic sublimes as arsenic(III) oxide leaving iron oxides, while roasting without air results in the production of metallic arsenic. Further purification from sulfur and other chalcogens is achieved by sublimation in vacuum or in a hydrogen atmosphere or by distillation from molten lead-arsenic mixture.

Applications

Agricultural

Roxarsone is a controversial arsenic compound used as a nutritional supplement for chickens

The toxicity of arsenic to insects, bacteria and fungi led to its use as a wood preservative. In the 1950s a process of treating wood with chromated copper arsenate (also known as CCA or Tanalith) was invented, and for decades this treatment was the most extensive industrial use of arsenic. An increased appreciation of the toxicity of arsenic resulted in a ban for the use of CCA in consumer products; the European Union and United States initiated this process in 2004. CCA remains in heavy use in other countries however, e.g. Malaysian rubber plantations.

Arsenic was also used in various agricultural insecticides, termination and poisons. For example, lead hydrogen arsenate was a common insecticide on fruit trees, but contact with the compound sometimes resulted in brain damage among those working the sprayers. In the second half of the 20th century, monosodium methyl arsenate (MSMA) and disodium methyl arsenate (DSMA) – less toxic organic forms of arsenic – have replaced lead arsenate in agriculture.

Arsenic is still added to animal food, in particular in the US as a method of disease prevention and growth stimulation. One example is roxarsone, which is used as a broiler starter by about 70% of the broiler growers since 1995. The Poison-Free Poultry Act of 2009 proposes to ban the use of roxarsone in industrial swine and poultry production. Alpharma, a subsidiary of Pfizer Inc., which produces Roxarsone, has voluntarily suspended sales of the drug in response to studies showing elevated levels of arsenic in treated chickens.

 

Medical use

During the 18th, 19th, and 20th centuries, a number of arsenic compounds have been used as medicines, including arsphenamine (by Paul Ehrlich) and arsenic trioxide (by Thomas Fowler). Arsphenamine as well as neosalvarsan was indicated for syphilis and trypanosomiasis, but has been superseded by modern antibiotics. Arsenic trioxide has been used in a variety of ways over the past 500 years, but most commonly in the treatment of cancer. The US Food and Drug Administration in 2000 approved this compound for the treatment of patients with acute promyelocytic leukemia that is resistant to ATRA. It was also used as Fowler’s solution in psoriasis. Recently new research has been done in locating tumors using arsenic-74 (a positron emitter). The advantages of using this isotope instead of the previously used iodine-124 is that the signal in the PET scan is clearer as the body tends to transport iodine to the thyroid gland producing a lot of noise.

In subtoxic doses, soluble arsenic compounds act as stimulants, and were once popular in small doses as medicine by people in the mid-18th century.

Name:Antimony                               Symbol:Sb

Type:Metalloid, Nitrogen group       Atomic weight:121.75

Density ρ 293 K:6.684 g/cm³           Atomic volume:18.22 cm³/mol

Discovered

The presence of antimony in historical artifacts indicates it was known to ancient civilizations. Combined with sulfur in stibnite (Sb₂S₃) it was used in Egyptian cosmetics four or five thousand years ago, as a black eyeliner.⁽¹⁾  

It’s likely that Roman author Pliny gave it the name stibium, from which the modern element symbol Sb was taken, in the first century AD. Stibnite is found most commonly, Pliny says, among the ores of silver. Pliny described stibnite’s use as a medicine. He also noted how if too strongly heated, it would turn to lead. What we understand now by this is the lead described by Pliny is actually the element antimony.

In the first half of the 1500s, Vannoccio Biringuccio wrote a description “Concerning Antimony and Its Ore”. This is an alchemical work. Biringuccio describes antimony sulfide as either “a monstrosity among metals” or, alternatively, “a material that is about to reach metallic perfection, but is hindered from doing so by being mined too soon”. He also warns against heating the antimony sulfide too strongly because this will produce a substance that “although this is very white and almost more shining than silver, it is much more brittle than glass.” This is a clear description of the element antimony.

Nicolas Lémery wrote his Treatise on Antimony in 1707. This was still not chemistry as we know it. In his writings, Lémery describes how acids prick the tongue because they contain spiky particles, while metals dissolve in acids because the sharp points of acids tear the metal particles apart.

The name “antimony” is derived from two Greek words: ‘anti’ and ‘monos’ which meaot alone. This results from the fact that antimony is infrequently found native, but usually combined with sulfur or with heavier metals such as copper, lead and silver.

States

Energies

Oxidation & Electrons

Appearance & Characteristics

Harmful effects

Like arsenic, which sits directly above it in the periodic table, the toxicity of antimony and its compounds varies according to the chemical state of the element. Many of the salts are carcinogenic.  

The metallic form is considered to be less active whereas stibine (SbH₃) and antimony trioxide are extremely toxic. Antimony is toxic and immediately dangerous to life or health at 50 mg m^-3 or above.  

Exposure to 9 milligrams per cubic meter of air (mg/m³) of antimony as stibnite for a long time can irritate your eyes, skin, and lungs. Breathing 2 mg/m³ of antimony for a long time can cause problems with the lungs (pneumoconiosis) heart problems (altered electrocardiograms), stomach pain, diarrhoea, vomiting and stomach ulcers. People who drank over 19 ppm of antimony once, vomited.

Characteristics

Antimony is metalloid, so it has some metallic properties but not enough to be classified as a true metal. Physically, it behaves like sulfur while chemically it is more metallic. Antimony’s electrical and thermal conductivity are lower than most metals’ conductivities. Antimony is a brittle, fusible, crystalline solid. It is easily powdered. Antimony also has the unusual property that (like water) it expands as it freezes. Four other elements expand when they freeze; silicon, bismuth, gallium and germanium. In addition to the usual form of antimony, there are two allotropes: yellow crystalline and amorphous black.

Uses

The major use of antimony is in lead alloys – mainly for use in batteries – adding hardness and smoothness of finish. The higher the proportion of antimony in the alloy, the harder and more brittle it will be. Alloys made with antimony expand on cooling, retaining the finer details of molds. Antimony alloys are therefore used in making typefaces for clear, sharp printing.  

Babbit metals, used for machinery bearings, are alloys of lead, tin, copper and antimony. These metals are hard but slippery and so ideal for use as bearings.

Antimony is used in the semiconductor industry as a-type dopant for silicon.  

Antimony trioxide is used as a flame retardant in adhesives, plastics, rubber and textiles.

Reactions

Compounds

Radius

Conductivity

Abundance & Isotopes

Abundance earth’s crust:0.2 parts per million by weight, 0.03 parts per million by moles. Abundance solar system:950 parts per billion by weight, 10 parts per trillion by moles.

Source

Most antimony is produced from stibnite (antimony sulfide, Sb₂S₃). It is also extracted as a byproduct of copper, gold and silver production.

Isotopes

31 whose half-lives are known, mass numbers 104 to 136. Of these, two are stable and found naturally in the percentages shown: ¹²¹Sb (57.36%) and ¹²³Sb (42.64%).

 

TESTS FOR GROUP V ELEMENTS

Phosphorus

Prolonged oxidation of any phosphorus compound, followed by standing in water, converts it to phosphate(V). This can then be detected by the formation of a yellow precipitate when heated with ammonium molybdate and nitric acid. Specific tests for various oxophosphates are known.

Arsenic

Because of its toxicity, it is ofteecessary to be able to detect arsenic when present only in small amounts in other substances. Arsenic present only in traces (in any form) can be detected by reducing it to arsine and then applying tests for the latter. In Marsh’s test, dilute sulphuric acid is added dropwise through a thistle funnel to some arsenic-free zinc in a flask; hydrogen is evolved and led out of the flask by a horizontal delivery tube. The arseniccontaining compound is then added to the zinc-acid solution, and the delivery tube heated in the middle. If arsenic is present, it is reduced to arsine by the zinc-acid reaction, for example :

AsO₄³  + 4Zn + 11H⁺  → AsH₃ + 4Zn²⁺ + 4H₂O

The evolved arsine is decomposed to arsenic and hydrogen at the heated zone of the delivery tube; hence arsenic deposits as a shiny black mirror beyond the heated zone.

Antimony and bismuth

As can be expected, antimony compounds resemble those of arsenic. In the Marsh test, antimony compounds again give a black deposit which, unlike that formed by arsenic compounds, is insoluble in sodium chlorate(I) solution. Solutions of many antimony and bismuth salts hydrolyse when diluted; the cationic species then present will usually form a precipitate with any anion present. Addition of the appropriate acid suppresses the hydrolysis, reverses the reaction and the precipitate dissolves. This reaction indicates the presence of a bismuth or an antimony salt. When hydrogen sulphide is bubbled into an acidic solution of an antimony or a bismuth salt an orange precipitate, Sb₂S₃, or a brown precipitate, Bi₂S₃, is obtained. Bismuth(III) sulphide, unlike antimony(III) sulphide, is insoluble in lithium hydroxide.

 

References:

1. The abstract of the lecture.

2. intranet.tdmu.edu.ua/auth.php

3. Atkins P. W. Physical chemistry / P.W. Atkins. – New York, 1994. – P.299‑307.

4. Cotton F. A. Chemical Applications of Group Theory / F. A. Cotton. ‑ John Wiley & Sons : New York, 1990.

5. Girolami G. S. Synthesis and Technique in Inorganic Chemistry / G. S. Girolami, T. B. Rauchfuss, R. J. Angelici. ‑ University Science Books : Mill Valley, CA, 1999.

6. Russell J. B. General chemistry / J B. Russell. New York.1992. – P. 550‑599.

7. Lawrence D. D. Analytical chemistry / D. D. Lawrence. –New York, 1992. – P. 218–224.

8. http://www.lsbu.ac.uk/water/ionish.html

9. http://en.wikipedia.org/wiki

The following website shows the reaction of VA group elements. It’s cool stuff! Check it out:

www.youtube.com/watch?v=odg2UTxgFHY

www.youtube.com/watch?v=2mU2LcWlWwM

www.youtube.com/watch?v=O1m6CP5XfDQ

www.youtube.com/watch?v=VuBpavnvH-c

 

Prepared by PhD Falfushynska H.